Synthesis and characterization of manganese oxides employed in VOCs abatement

Synthesis and characterization of manganese oxides employed

in VOCs abatement

Luciano Lamaita, Miguel A. Peluso, Jorge E. Sambeth *, Horacio J. Thomas

Centro de Investigacion y Desarrollo en Ciencias Aplicadas (CINDECA) ‘‘Dr. Jorge J. Ronco’’—UNLP,

CONICET—47 No 257 (B1900AJK) La Plata, Buenos Aires, Argentina

Received 10 October 2004; received in revised form 24 March 2005; accepted 30 March 2005

Available online 25 May 2005

Abstract

Two MnO2 were prepared by different methods: (i) MnCO3 decomposition; and (ii) oxidation of acid MnSO4 solution. The combined use

of X-ray diffraction (XRD), DRIFTS, DRS–UV–vis and TGA techniques revealed: (i) the presence of Mn4+–Mn3+; (ii) a poor crystallinity;

and (iii) OH species at temperatures higher than 200 8C.

The catalysts were tested in the complete oxidation of ethanol to carbon dioxide and water, and the reaction results showed that the

catalytic activity of both samples were larger than MnO2 (pyrolusite, b-MnO2) and Mn2O3 (like bixbyite) oxides. The high activity of the

prepared oxides can be related to the simultaneous presence of Mn3+ and Mn4+ ions and OH� species generated by Mn4+ vacancies. In

addition, in situ DRIFT spectroscopy demonstrated a different mechanism of adsorption-reaction of ethanol between the two synthesized

manganese oxides.

# 2005 Elsevier B.V. All rights reserved.

Keywords: g-MnO2 nsutite; Manganese oxides; DRIFTS; Ethanol oxidation; VOCs

www.elsevier.com/locate/apcatb

Applied Catalysis B: Environmental 61 (2005) 114–119

1. Introduction

Among the transition metal dioxides, MnO2 probably

exhibits the largest number of polymorph structures.

Pyrolusite (b-MnO2) is the most stable polymorph of

manganese dioxide and has the structure of rutile [1].

Ramsdellite (R-MnO2) is closely related to rutile except for

the fact that double chains replace the single chains of

edge—sharing octahedra [2] and nsutite or g-MnO2 has a

highly disordered structure and has been described as an

intergrowth of elements of pyrolusite and ramsdelllite [3].

Manganese dioxides can be classified according to the

number of MnO6 units and the number of MnO6 octahedral

chains between two basal layers to form tunnel openings.

They are usually symbolized T(m, n) where n and m stand

for the dimension of the tunnels in the two directions

perpendicular to the chains of edge-sharing octahedra [4].

Thus, pyrolusite is T(1, 1) and ramsdellite is T(1, 2), so that

* Corresponding author. Fax: +54 221 425 4277.

E-mail address: [email protected] (J.E. Sambeth).

0926-3373/$ – see front matter # 2005 Elsevier B.V. All rights reserved.

doi:10.1016/j.apcatb.2005.03.014

g-MnO2 could be represented as T(1, 1)–T(1, 2) inter-

growths.

Among the various polymorphs of MnO2 reported in the

literature (Chabre and Pannetier [4] have reported about 14

modifications) the g-MnO2 (nsutite) is the best-known

MnO2 polymorph used by the battery industry. g-MnO2 has

been extensively investigated for its important application as

cathode material for dry cells. According to McLean et al.

[5], structural and chemical defects are responsible for

electrochemical properties by the H+ insertion and increas-

ing the Fermi Level energy. The Volkenshtein Electronic

Theory [6] states that the catalytic activity is a function of

the electronic properties of solids, so the g-MnO2 phase

could be an excellent catalyst.

In this sense, Lahousse et al. [7] and Peluso et al. [8] have

shown that the g-MnO2 and one sample of Mn3+/g-MnO2,

respectively, are very suitable for VOC’s removal catalyst.

The aims of this paper are: (1) to characterize g-MnO2

catalysts prepared by different methods of synthesis; (2) to

study the catalytic activity in the complete oxidation of

ethanol to carbon dioxide and water, and to correlate this

L. Lamaita et al. / Applied Catalysis B: Environmental 61 (2005) 114–119 115

activity to the structural properties of the catalysts; and (3) to

propose a possible adsorption–oxidation mechanism.

2. Experimental

2.1. Sample preparation

Four manganese oxide samples were used in this work.

Two of them were prepared in order to obtain the g-MnO2

phase. In a first attempt to obtain this phase, MnCO3 (Riedel

de Haen RG) was oxidized at 500 8C under flowing oxygen

(50 cm3/min) with a heat rate of 20 8C min�1 [8]. This oxide

is named hereinafter M1.

The second manganese oxide (M2) used in this work was

prepared by the Netto et al. procedure [9]. A solution of

MnSO4�H2O dissolved in 3 M H2SO4 was oxidized with

flowing oxygen (30 cm3/min). pH was raised up to 3 with

NaOH. After being filtered and washed with deionised

water, the oxide was dried at 100 8C for 24 h.

For the purpose of comparison, a commercial reagent

grade MnO2 (Baker 99.9%), named M3, has been

investigated. Finally, a synthetic pyrolusite (M4) prepared

by thermal decomposition of solid Mn(NO3)2�4H2O was

also studied [10].

The list of the oxides, together with the relative notation,

is presented in Table 1.

2.2. Characterization

The structure of the studied samples was characterized

through X-ray diffraction (XRD) by using a Phillips PW

1390 instrument with Ni-filtered Cu Ka radiation

(l = 1.540589 A). The diffraction patterns were taken at

room temperature in the range 5 < 2u < 708. For phase

identification purposes, the JCPDS database of reference

compounds was used [11].

The BET specific surface areas were measured by N2

adsorption at the liquid nitrogen temperature (77 K) in a

Micromeritics Accussorb 2100 D sorptometer.

DRIFTS spectra were obtained in a Bruker IFS66

infrared spectrometer with a KBr optics and DTGS detector

in the 600–4000 cm�1 range. An environmental DRIFTS

chamber equipped with KBr window (Spectra Tech 0030-

103), allowing in situ treatments up to 300 8C and 1 atm was

coupled to the spectrometer. The spectra were obtained by

co-adding 200 scans collected at 4 cm�1.

Table 1

BET specific surface areas of the manganese oxides prepared in this study

Sample MnOx precursor BET specific area (m2 g�1)

M1 MnCO3 11

M2 MnSO4�H2O 55

M3a – 0.1

M4 Mn(NO3)2�4H2O 6.1a MnO2 baker 99%.

The DRS–UV–vis spectra were recorded on a Cary 2300

spectrophotometer in the region 200–2000 nm.

Thermal analyses (TG) of the catalysts were performed

on a DTA-TG 50 Shimadzu analyzer between room

temperature and 900 8C; small portions (14–16 mg) of the

samples were used at 10 8C min�1 in a flowing atmosphere

of air (30 cm3 min�1).

2.3. Catalytic activity

Ethanol was chosen as a probe molecule because it is an

important VOC produced in fermentation and sponge dough

processes and the industrial production of beer, food or

pharmaceutical products [12].

Catalytic test was carried out at atmospheric pressure in a

continuous flow tubular glass reactor. Two hundred

milligrams of the sample were loaded in form of fine

powder. The total gas flow of ethanol–air was 20 cm3 min�1

and the work temperatures were between 100 and 300 8C.

The feed composition was 900 ppm of ethanol in air, with an

excess of O2.

The reactants and the reaction products of a possible

incomplete reaction were analyzed by using an on-line gas

chromatograph (Shimadzu GC 8A, Porapak T for the

analysis of ethanol, CO2, acetaldehyde and a 5A zeolites for

the analysis of CO) and quantified by using a TCD. For the

purpose of comparison, the temperature at which the

conversion of the organic compound reaches 50% (T50) was

used to analyze the activity of the catalysts in VOCs

oxidation. The low T50 shows the best catalytic activity.

The adsorbed species formed during the reaction on the

manganese oxides were studied by in situ DRIFT spectro-

scopy by using the same experimental conditions achieved

in the studies of catalytic activity.

3. Results and discussion

X-ray diffraction was used to determine the bulk

crystalline phases in the samples. The diffraction patterns

of the manganese oxides are shown in Fig. 1. The XRD

Fig. 1. X-ray powder diffraction patterns of the manganese oxides.

L. Lamaita et al. / Applied Catalysis B: Environmental 61 (2005) 114–119116

Table 3

IR frequency vibration of manganese oxides

Sample IR bands (cm�1) References

M1 727 633 544

M2 718 633 542

M3 – 616 560

M4 – 615 560

b-MnO2 618 562 [19,20]

R-MnO2 743 649, 630 522 [19,20]

R-MnO2 740 687 529 [19]

g-MnO2 705 545 [19]

patterns of the commercial MnO2 sample (M3) show the

typical spectrum of b-MnO2 phase (pyrolusite) (JCPDS, 24-

0735). The oxide prepared by decomposition of

Mn(NO3)2�4H2O (M4) also shows the spectrum of

pyrolusite.

A more complex XRD pattern is obtained for M1 sample.

The signal at 2u 338 indicates the presence of a-Mn2O3,

bixbyite phase. Additionally, the broad and poorly resolved

peak in the 2u range 37–398 could indicate the presence of a

mixture of Mn(IV) and Mn(III) oxides, according to Arena

et al. [13]. Thus, this oxide is composed by a mixture of

Mn2O3 and an amorphous Mn4+–Mn3+ oxide.

Chabre and Pannetier [4] have reported that the XRD

spectrum of the g-MnO2 (nsutite) phase is always of rather

poor quality and consists, at best, of a small number of sharp

and broad lines on top of a diffuse background. The X-ray

powder diffraction pattern of M2 shows small and broad

lines at 2u 228, 378, 388 and 568, which reveals the presence

of the g-MnO2, according to database of the reference

compound (JCPDS, 17-0510).

The BET specific surface areas of the samples are

summarized in Table 1. The results showed a highest surface

area for M2 sample (nsutite phase oxide). Its high surface

area could be explained by a higher number of point defects

on this oxide surface [14]. As it can be seen from the table,

the M1 sample surface area is higher than that of M3 and

M4, but lower than M2. Finally, although M4 and M3

present the pyrolusite structure, M4 has a 60 times higher

surface area than M3.

The DRS–UV–vis position bands of different manganese

oxides can be seen in Table 2. For the M3 and M4 samples,

the spectra are characterized by an absorption band near

340 nm, which is assigned to Mn4+ [15]. The M1 and M2

samples also present that absorption. On the other hand, the

spectrum of M1 and M2 samples presents an absorption

band between 450 and 550 nm, which could be due to charge

transfer transitions of octahedral Mn3+ [16]. These features

suggest the existence of Mn4+ and Mn3+ in the bulk of the

M1 and M2 oxides, in agreement with XRD data.

The results of DRIFT spectroscopy of the studied

manganese oxides are reported in Table 3. Neither CO32�

nor SO42� groups were detected by DRIFTS spectroscopy in

the fresh catalysts.

IR spectroscopy has proven to be useful for the

identification of the nsutite phases [17], and yields more

reliable information than the X-ray diffraction technique. The

IR bands in the region 1000–400 cm�1 reveal information

Table 2

DRS–UV–vis bands of manganese oxides

Sample Mn4+ (nm) Mn3+ (nm)

M1 450–550 340

M2 450–550 340

M3 450–550

M4 450–550

Mn2O3 340

about MnO6 octahedra [18]. Moreover, Potter and Rosmann

[19] stated that the infrared spectroscopy is sensitive to

octahedral polymerization. Besides, Julien et al. [20]

observed a general decrease in band wavenumber with

increasing octahedral polymerization. These author assigned

the bands at 545 cm�1 in pyrolusite and at 522 cm�1 in

ramsdellite to the vibration due to the displacement of the

oxygen anions relative to the manganese ions along the

direction of the octahedral chains.

The M1 bands at 727 and 544 cm�1 could be assigned to

g-MnO2 and the band at 633 cm�1 to Mn2O3, respectively.

In the case of M2, DRIFTS bands are in accordance with

those reported by Potter and Rossman [19] for the g-MnO2

phase. In addition, these authors have also reported that the

characteristic bands of ramsdellite are at 522, 630, 649 and

743 cm�1. Taking into account the correlation of infrared

spectra made by Potter and Rossman [19] and Julien et al.

[20], it is possible to determine the structural characteristics

of our manganese oxides (M1, M2). In Fig. 2, the positions

of the IR bands are shown as a function of edge shared

octahedra. Both M1 and M2 samples are placed between the

ramsdellite and the pyrolusite, in agreement with the fact

that the g-MnO2 is a mixture of ramsdellite and pyrolusite

structures [3].

Fig. 2. Frequency positions of the IR bands, a function of the average MnO6

octahedral polymerization in manganese oxides. (ramsdellite bands, Potter

and Rosmann [18]).


According to the results of the characterization techni-

ques we can say that:

(a) M

Fi

1 catalysts are constituted by an amorphous mixed

Mn2O3–g-MnO2 oxide.

(b) M
2 is an oxide of the g-MnO2 (nsutite) type with
manganese(III) and manganese(IV) cations.

(c) M
3 and M4 catalysts have the b-MnO2 (pyrolusite)
structure and the manganese oxidation state is IV in both

catalysts.

Fig. 4. Total ethanol conversion as function of temperature.

Kanungo et al. [21] have reported that the thermal

analysis of manganese oxides between 25 and 900 8C in air

presents three different thermal effects: (a) loss of physically

adsorbed molecular water in the temperature range

100–200 8C; (b) loss of chemically bound water and a

change of crystalline phase between 200 and 400 8C; and (c)

the transformation of MnO2 to Mn2O3 around 550 8C.

Likewise, Vileno et al. [22] in their studies of OMS

manganese oxides also reported three thermal events: loss of

physisorbed water from 100 to 130 8C; loss of chemisorbed

water, inside or outside the tunnels, at 200 8C; and loss of

oxygen.

Fig. 3 depicts TG curves for the studied oxides except for

M4. As it can be seen, decomposition of MnO2 to Mn2O3 in

M1 sample occurs at 580 8C. The other catalysts show a

completely different behaviour. In fact, M2 sample is still

losing water even at 800 8C. This is in agreement with Petit

et al. [23] who pointed out that the g-MnO2 is non-

stoichiometric unless the temperature is raised up to 800 8C.

This is due to the presence of OH-groups in the bulk of the

oxide. They have also shown that this oxide presents

electrochemical activity at temperatures higher than 450 8C.

Finally, M3 is the most stable oxide. There is not physically

or chemically bound water, since no weight loss is present

until the temperature reaches 610 8C. At this stage, the

transformation to Mn2O3 occurs with approximately 10%

weight loss, and it is in agreement with the calculated value

(9.2%).

g. 3. TGA curves for manganese oxides: (a) M1; (b) M2; (c) M3.

The catalytic performances for ethanol oxidative

destruction on manganese oxide catalysts are shown as

function of temperature in Fig. 4. In all the samples tested,

CO2, H2O and acetaldehyde were formed during ethanol

oxidation, and no other partial oxidation products were

detected. Preliminary experiments showed that the com-

mercial manganese carbonate did not exhibit any activity at

all for the catalytic oxidation of ethanol.

Fig. 4 shows that the oxide with the nsutite structure (M2

sample) is a more active catalyst with a T50 more than 40 8Clower than that of the catalyst prepared by decomposition of

MnCO3 (M1 sample). Total combustion occurs at 200 8C for

the M2 catalyst, and at around 215 8C for M1 catalysts. In

both catalysts (M1 and M2), the selectivity to acetaldehyde

was at a peak when the conversion was low (less than 20%).

It was also observed that carbon dioxide was the only

product, when the conversion reached 30%. Both pyrolusite

type catalysts (M3 and M4) have the lowest activity of the

catalysts with a T50 close to 250 8C. The combustion is not

total until the temperature reaches 300 8C.

As we have seen, the most crystalline and lowest specific

surface area oxide (M3, b-MnO2) has the lowest activity,

and the most active catalyst (M2) has the lowest crystallinity

and the highest specific surface area. These results are in

agreement with Stobbe et al. [24] and Shaheen and Selim

[25], who have reported that the oxides with high degree of

crystallinity had a pronounced decrease in catalytic activity.

Nevertheless, not only the surface area but also the

structure of the oxides influences in the catalytic activity of

this series of oxides. M4 sample has a 60 times higher

specific surface area than that of M3 sample (Table 1), but

the activity of the former is lower than that of the latter (see

Fig. 4).

Besides the specific surface area and the crystallinity of

the catalysts, the existence of the Mn3+–Mn4+ couple [26–

28], the presence of cationic vacancies and OH groups [29]

are responsible for the high activity achieved by the nsutite-

phase oxide (M2).

L. Lamaita et al. / Applied Catalysis B: Environmental 61 (2005) 114–119118

As far as the oxidation state is concerned, the UV–vis and

DRIFTS spectra showed that M1 and M2 samples contained

Mn3+ and Mn4+ ions, whereas pure MnO2 without any

amount of Mn3+ constituted M3 and M4 samples. Although

both M1 and M2 samples contain Mn3+ in their structure,

they have different crystalline structures. In M1 sample,

Mn3+ is present in the form of crystalline Mn2O3 and it is

also present in a mixed Mn4+–Mn3+ amorphous oxide, where

the presence of the Mn4+ ions could be associated with g-

MnO2 phase. But in the case of M2 sample, Mn3+ cations

could be distributed in a nsutite matrix replacing Mn4+

cations. In this oxide, no Mn2O3 is detected.

Taking into account the catalytic activity results, oxides

with Mn valence between 3 and 4 were more active than

those with pure Mn (+4). M1 catalyst had a mixed structure

of Mn2O3 and g-MnO2, as it was confirmed by DRIFTS

spectra. Therefore, the presence of the couple Mn3+–Mn4+

could be responsible for the best catalytic performances

compared to b MnO2 and synthetic Mn2O3 [8].

The presence of Mn4+ vacancies in the nsutite structure is

followed by the formation of OH groups, in order to

equilibrate charges. The TGA results showed OH species at

200 8C when conversion was complete for M2. So, we can

say that OH groups played an important role in the oxidation

mechanism of VOC’s.

Fig. 5 shows the DRIFTS spectra of adsorbed species over

M1 catalysts. The spectrum of ethanol adsorption over M1

sample from RT to 70 8C shows bands assigned to ethoxide

species (1640–1650 cm�1 enolic anion; 1420 cm�1 dOH; near

1110 cm�1 C–O stretching and 1050 cm�1 ethoxy group)

[30,31]. At 70 8C, a new band at 1746 cm�1 attributed to

vibration nC O (acetaldehyde) is detected. At 150 8C, bands

begin to be formed at 1430 and 1555 cm�1, that can be

assigned to a chelate bidentate acetate [29]. Above 150 8C,

three bands are detected, two of them (1425–1570 and

Fig. 5. In situ DRIFTS spectra of the adsorption of ethanol on M1 catalyst

as a function of the temperature: (a) RT; (b) 70 8C; (c) 100 8C; (d) 120 8C;

(e) 150 8C; (f) 200 8C.

1045 cm�1) assigned to bridged bidentate acetate and at

1350 cm�1, a third band attributed to the monodentate acetate

species. The fact that different organic species are found in

M1 enables us to suppose that there are two different

mechanisms of adsorption-reaction, in good agreement with

Batiot and Hodnett [32] and Blasim-Aube et al. [33], who

have postulated that the ethanol oxidation to carbon dioxide

over catalyst can be described as:

Then, the formation of different adsorbed species

could depend upon the different bond dissociation enthalpy

of the weakest C–C and C–H bonds that ethanol and

acetaldehyde (389 kJ/mol C–H ethanol, 358 and 349 kJ/mol

C–C and C–H acetaldehyde) present when they are oxidized

to CO2.

Fig. 6 shows the infrared spectra of the adsorbed species

over M2 as a function of the reaction temperature. The

spectrum is different from the one recorded over M1 sample.

At room temperature, bands are detected at: (i) 1290 cm�1

attributed to adsorbed ethanol; (ii) 1023 cm�1 assigned to

ethoxy species; and (iii) 1475 cm�1 corresponding to a

vibration dCH2. In the range 70–200 8C, the main bands are

detected at: (i) 1270–1500 cm�1 attributed to monodentade

carbonate; and (ii) 1736 cm�1 assigned to carbonyl species

(acetaldehyde). In addition, the shoulder at 1177 cm�1 could

be due to ethoxy species coordinated over metal sites.

These results showed that the adsorption-reaction path-

way of ethanol over M2 was produced by the presence of

different adsorbed species with respect to M1 catalyst.

Taking into account these results, we propose the

following pathways for ethanol oxidation:

Fig. 6. In situ DRIFTS spectra of the adsorption of ethanol on M2 catalyst

as a function of temperature: (a) RT; (b) 70 8C; (c) 100 8C; (d) 150 8C; (e)

200 8C.


The different mechanism could be attributed to the

formed solid. According to Ruetschi [34], g-MnO2 contains

Mn vacancies compensated by four protons that are attached

to the vacancy (OH formation). So, the interaction between

ethanol and oxygen and/or hydroxyl centres on the surface

of M2 catalyst could favour the formation of monodentade

carboxylate species and improve the catalytic performances.

4. Conclusions

Two manganese oxides were synthesized in order to

obtain the nsutite phase (g-MnO2) and to make a comparison

with two dioxides with the pyrolusite structure (b-MnO2).

The oxides with the nsutite (g-MnO2) structure (M2 and M1)

are more active catalyst than the solids with pyrolusite

structure (b-MnO2, M3 and M4).

The characterization results allow us to conclude that the

high activity of both catalysts could be due to its poor

crystallinity, the existence of the couple Mn3+–Mn4+ and the

presence of Mn4+ vacancies.

The in situ DRIFT spectroscopy showed that the two

synthetic manganese oxides presented different mechanisms

of adsorption-reaction of ethanol.

The mechanism of adsorption and the changes in the

activity between M1 and M2 catalysts could be explained

when assuming that these catalysts contain different

structures: Mn2O3/g-MnO2 and g-MnO2, respectively.

Acknowledgments

The authors acknowledge Prof. L. Cadus (UNSL,

Argentina) for TGA analysis, and Prof. G. Minnelli

(U. La Sapienza, Italy) for the DRS–UV–vis spectra. M.

Peluso thanks Comision de Investigaciones Cientıficas of

Buenos Aires (Argentina) for his Ph.D. fellowship.

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Synthesis and characterization of manganese oxides employed in VOCs abatement

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